Redox and Electrochemistry (BLB chapter 20, p.723)

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1 Redox and Electrochemistry (BLB chapter 20, p.723) Redox is short for reduction/oxidation Redox chemistry deals with changes in the oxidation states of atoms Oxidation States All atoms have an oxidation state (or oxidation number). Atoms that are unbonded, or bonded only to atoms of the same element have a zero oxidation state. e.g. He, O 2, H 2 - the oxidation state of each atom is zero. The oxidation states of all atoms in a molecule add up to zero e.g. H 2 O - each H has an oxidation state = +1 (total of +2) and the O has an oxidation state of -2 The sum of oxidation states of all atoms in H 2 O is zero The oxidation states of all atoms in a polyatomic ion add up to the charge on the ion e.g. CO each O has an oxidation state of -2 (total of -6) and the C has an oxidation state of +4 Sum of oxidation states of all atoms in CO 2-3 is -2 The oxidation states of most elements vary depending on the molecule they are in. However, many elements spend most of their time in one particular oxidation state (when bonded). These oxidation states are closely related to the element s group. Element Oxidation State Group 1 (Alkali metals) Li, Na, etc. +1 Group 2 (Alkaline earths) Be, Mg, etc. +2 Group 13 B, Al, etc. +3 Oxygen -2 Group 17 (halogens) F, Cl, Br, I -1 Transition metals can have several different oxidation states The oxidation state of a transition metal in a particular compound is indicated by a Roman numeral (I for 1, II for 2, etc.) e.g. iron (II) carbonate, FeCO 3 and iron (III) chloride, FCl 3 1

2 Redox Reactions In a redox reaction, electrons are transferred between atoms (as well as a rearrangement of atoms, like all reactions) If an atom loses an electron, its oxidation state increases by 1 it is oxidized If an atom gains an electron, its oxidation state decreases by 1 it is reduced (Forgive me for the next bit.) You can remember this using the mnemonic device; LEO the lion says GER LEO: you Lose Electrons in Oxidation GER: you Gain Electrons in Reduction e.g. adding iron to hydrochloric acid gives a redox reaction Fe + 2 HCl FeCl 2 + H 2 o On the reactant side, oxidation states are; Fe = 0, H = +1, Cl = -1 o On the product side, oxidation states are; Fe = +2, H = 0, Cl = -1 o So, Fe is oxidized its oxidation state gets more positive (0 to +2) o And, H is reduced its oxidation state gets more negative (+1 to 0) o o o In every redox reaction, something is oxidized and something is reduced The thing that is oxidized is called the reducing agent (or reductant) The thing that is reduced is called the oxidizing agent (or oxidant) Redox reactions can be written as two half reactions One half reaction shows reduction One half reaction shows oxidation e.g. for the reaction given above; Fe Fe e - Oxidation 2 H e - H 2 Reduction The Cl- ion is not included in either half-reaction because its oxidation state doesn t change during the reaction, even though it switches from H to Fe. NOTE: the number of moles of electrons transferred in a redox reaction is simply the coefficient of e - in one half reaction. e.g. MnO Cr + 2 H 2 O MnO 2 + Cr OH - Half reactions: Mn e - Mn 4+ Cr Cr e - There are 3 moles of electrons transferred in this reaction. 2

3 Reduction Potentials (E 0 ) All half reactions, like the ones above (p.2), have a potential (also called a voltage). The potential tells us how happy the reaction is if it occurs A positive potential means it s happy the half reaction is spontaneous A negative potential means it s unhappy the half reaction is not spontaneous Potential is also called voltage, because the half reaction will produce this voltage if it is part of a Galvanic cell. The standard reduction potential, E 0 (given in units of V, Volts) of every half reaction is given on the AP test, and also in a table on pages of BLB. Standard conditions: 25 C and all solutions are 1 M concentration The reduction potential is the potential for the reduction of an element or ion (electrons are added, remember LEO says GER; Gain Electrons in Reduction) e.g. Li + + e - Li (s) E 0 = V e.g. F 2 (g) + 2 e - 2 F - E 0 = 2.87 V Only the reduction potentials are given For oxidation, the reduction half reaction is reversed and the sign of the potential is changed (+ to -, or to +) e.g. Li (s) Li + + e - E 0 = 3.05 V e.g. 2 F - F 2 (g) + 2 e - E 0 = V The more positive the reduction potential, the more likely it is that the half reaction will take place. The more negative the reduction potential, the less likely it is that the half reaction will take place. Half reactions with a negative potential are not spontaneous (they don t happen on their own), but they will happen if the full reaction has a positive potential. e.g. (a) Is solid iron oxidized in a solution of Zn 2+ ions? (b) Is solid zinc oxidized in a solution of Fe 2+ ions? The reduction potentials for iron and zinc are; Fe e - Fe (s) E 0 = V Zn e - Zn (s) E 0 = V Neither reduction process is spontaneous (both have negative potential) (a) The redox half reactions would be; Fe (s) Fe e - E 0 = 0.44 V Zn e - Zn (s) E 0 = V The total potential is simply the sum of the half reaction potentials. E 0 reaction = 0.44 V + (-0.76 V) = V Reaction is NOT spontaneous 3

4 (b) The redox half reactions would be; Fe e - Fe (s) E 0 = V Zn (s) Zn e - E 0 = 0.76 V E 0 reaction = V V = 0.32 V Reaction IS spontaneous (solid iron forms on surface of zinc) Non-spontaneous reduction of iron takes place because oxidation of zinc is very spontaneous (i.e. zinc oxidation has bigger +E 0 than E 0 of iron reduction). Voltage and Spontaneity Although we have already noted that a redox reaction is spontaneous if the potential for the reaction (E 0 reaction) is positive, we can show this formally by calculating G 0, the Gibbs free energy change of the reaction. G 0 < 0 (i.e. negative) for spontaneous processes G 0 > 0 (i.e. positive) for non-spontaneous processes G 0 = -nfe 0 G 0 is the standard Gibbs free enerfy change (kj / mol) n is the number of moles of electrons transferred in the reaction (mol) F is Faraday s constant, 96,500 coulombs / mol (i.e. the charge on 1 mole of electrons) E 0 is the standard reduction potential (V) Voltage and Equilibrium If we know the potential of a redox reaction, we can calculate the equilibrium constant in two ways; (1) calculate G 0 as shown above, then use, G 0 = -RT lnk (2) Calculate K directly using; E 0 = RT lnk nf R is the gas constant in J/mol.K (8.314), which is the same as V.coulomb/mol.K T is the absolute temperature (K) n and F are the same as the in previous equation NOTE: at 25 C this equation can be simplified to; log K = n E Remember that: if K > 1 the forward reaction is favored (E 0 is positive) if K < 1 the backward reaction is favored (E 0 is negative) 4

5 Galvanic (Voltaic) Cells In a Galvanic cell, the two half-reactions of a redox reaction are separated. The transferred electrons are forced to pass through a wire The electrons can do electrical work (e.g. light a bulb) An example Galvanic cell is shown here; What do you need to know? 1. What are the half reactions Ni KNO 3 Salt Bridge Zn taking place in each beaker? 2. Which electrode is the cathode and which is the anode? 3. Which way are the electrons moving along the wire? 4. What is the net ionic equation for the overall reaction, and 1.0 M Ni(NO 3 ) M Zn(NO 3 ) 2 what is the standard cell potential, E 0? 5. What does the salt bridge do (and what would happen if it was removed)? 6. What is the cell potential for non-standard concentrations of solution (i.e. not 1.0 M)? e.g. what if [Ni 2+ ] was changed to 0.10 M from 1.0 M? 1. To find the half reactions, look up Ni and Zn in the standard reduction table; Ni e - Ni (s) E 0 = V Zn e - Zn (s) E 0 = V Both potentials are negative, so neither metal really wants to be reduced. But the reduction of Ni 2+ has a more positive (less negative) reduction potential, so it is easier to reduce than Zn 2+. Therefore, Ni 2+ is reduced and Zn is oxidized, and the half reactions are; Ni e - Ni (s) E 0 = V Zn (s) Zn e - E 0 = 0.76 V 2. The electrodes are the pieces of metal sticking into the solutions (the wire is attached to them). One is called the cathode (it has a + sign in an electrical circuit), the other is called the anode (it has a sign in an electrical circuit). Oxidation occurs at the anode (use Mr. maddox to remember ANODE OXIDATION) Reduction occurs at the cathode (RED CAT to remember REDUCTION CATHODE) In the example above, Ni 2+ is reduced, so the electrode in contact with the Ni 2+ solution is the cathode this is the Ni electrode In the example above, Zn is oxidized, so the zinc electrode is the anode. 3. Although the current in an electrical circuit travels from cathode to anode (+ to -), the electrons themselves actually travel from the anode to the cathode (- to +). This makes sense if we remember that oxidation occurs at the anode, and oxidation produces electrons. These electrons then travel along the wire to the cathode, where they are used in the reduction process. In this cell, the electrons move along the wire from Zn to Ni. Bulb Wire 5

6 4. The net ionic equation is the sum of the two half-reactions (balanced, if necessary), and the cell potential is the sum of the two half-reaction potentials; Ni 2+ + Zn Ni + Zn 2+ E 0 = 0.76 V + (-0.25 V) = 0.51 V 5. Without the salt bridge, reduction of Ni 2+ to Ni would leave an excess of NO 3 - ions in the Ni(NO 3 ) 2 solution. The solution would become negatively charged and electrons traveling along the wire towards the Ni cathode would be repelled. Similarly, the oxidation of Zn to Zn 2+ would create a positively charge solution on the zinc side. The electrons would be attracted to the solution and not want to travel away from the Zn anode. Without the salt bridge, no current would flow (the bulb would not light up). The salt bridge provides ions that balance the charges formed by the oxidation and reduction processes (but do not react with anything in the cell). A semi-permeable connection that allows limited passage of ions between the two solutions can be used instead of the salt bridge. 6. The Nernst Equation is used to calculate cell potentials under non-standard conditions (concentrations 1.0 M). This equation is given on the AP equation sheet. E cell = E 0 cell - RT lnq nf At 25 C, it can be simplified to; E cell = E 0 cell log Q n E cell is the cell potential under the non-standard conditions E 0 cell is the cell potential under standard conditions (using standard reduction table) R is the gas constant, J/mol/K (this is the same as volt.coulomb/mol/k) T is the absolute temperature (K) n is the number of moles of electrons transferred (number of e - in one half reaction) F is the Faraday constant, 96,500 coulombs/mol Q is the reaction quotient (i.e. same as the equilibrium expression) For the cell given above, Ni 2+ + Zn Ni + Zn 2+, Q = [Zn 2+ ] [Ni 2+ ] Note that Zn and Ni are solid and do not appear in the Reaction quotient expression (or in an equilibrium expression) For the example above, E 0 cell = 0.51 V, n = 2, and Q = 1.0 M / 0.10 M = 10. So, E cell = 0.51 V - (0.0592) (log 10) / 2 = 0.48 V Note: Q = 1 for a standard cell. If Q > 1, the cell potential is less than the standard potential If Q < 1, the cell potential is greater than the standard potential 6

7 Additional Note on Electrolysis Non-spontaneous redox reactions can be forced to take place using a battery. It will only work if the battery pushes electrons in the opposite direction to the spontaneous redox reaction, and has a higher voltage. A common example is electrolysis of molten NaCl to form liquid Na and Cl 2 gas. Reduction at the cathode: 2 Na + + 2e - 2 Na (l) E 0 = V Oxidation at the anode: 2 Cl - Cl 2 (g) + 2e - E 0 = V Total standard cell potential; E 0 = V So, we need a battery with voltage greater than 4.07 V to electrolyze molten NaCl. Note: an NaCl solution will not work, because the water will be electrolyzed before the NaCl. Extra Note: Naming Ions and Compounds What You Really Need to Know You should be able to name all elements in the first four rows of the Periodic Table, given their symbols. Remember that group 1 elements (and H) like to form +1 ions Remember that group 2 elements like to form +2 ions Remember that group 16 elements (e.g. O, S) like to form -2 ions Remember that group 17 elements (e.g. Cl, Br) like to form -1 ions Transition metals usually make more than one type of ion, though all will be positive - use the standard reduction potential table to help, or guess +2, or +1 Molecules you should know H 2 O water SO 2 sulfur dioxide NH 3 ammonia HNO 3 nitric acid H 2 SO 4 sulfuric acid HCl hydrochloric acid HF hydrofluoric acid HI hydroiodic acid HBr hydrobromic acid CO carbon monoxide CO 2 carbon dioxide Cl 2, F 2, Br 2, I 2, H 2, N 2, O 2 elements are diatomic Learn the first five alkanes, plus alcohol, amine, and carboxylic acid versions Ions you should know O 2- - oxide PO phosphate S 2- - sulfide HCO hydrogen carbonate (or bicarbonate) SO sulfate CrO chromate NO nitrate Cr 2 O dichromate (good oxidizer) CO carbonate MnO permanganate (good oxidizer) OH - - hydroxide NH ammonium CN - - cyanide SCN - - thiocyanate CH 3 CO acetate (acetic acid without an H + - acetic acid is also called ethanoic acid) 7

Electrochemistry. Galvanic Cell. Page 1. Applications of Redox

Electrochemistry. Galvanic Cell. Page 1. Applications of Redox Electrochemistry Applications of Redox Review Oxidation reduction reactions involve a transfer of electrons. OIL- RIG Oxidation Involves Loss Reduction Involves Gain LEO-GER Lose Electrons Oxidation Gain